Valency: The combining power of an atom, and is equal to the number of hydrogen atoms it could combine with or displace.
Number of valence electrons refer to the number of electrons in an atom's Valence (outer) shell.
e.g. Nitrogen has a valency of -3, but has 5 valence electrons.
How the periodic table works
The periodic table is ordered in a way that reflects its electronic structure.
Period number represents the number of electron shells of an element.
e.g. Calcium has an electron configuration (standard notation) of 2, 8, 8, 2. From this, we know that Calcium has 4 electron shells. Thus, Calcium has a period number of 4.
Group number represents the number of valence electrons of an element.
There are exceptions: e.g. the transition metals, however for yr 10 you don't need to know about why that happens.
For example, Chlorine has a group number of 17, however if we pretend the transition metals don't exist, it would have a group number of 7. Thus, Chlorine would have 7 valence electrons. Chlorine's electron configuration is 2, 8, 7 and does indeed have 7 valence electrons.
3 (main) types of bonding:
But why do elements bond? In general, atoms want to each a stable electron configuration. What that means is that they want their outer (valence) shell to be full of electrons (in general this is 8). They can do this in a variety of ways, by losing, gaining or sharing electrons.
We can also say they want the electron configuration of a noble gas, as noble gasses already have this configuration.
Note that bonding works on a spectrum of sorts. We can have pure metallic on one side, pure ionic in the middle, and pure covalent on the other side. However, there will always be certain compounds/molecules that do not fit in this category. E.g. $SiO_2$ can be considered 'covalent', however due to Silicon being a metalloid, $SiO_2$ (silica) is actually a covalent network substance, unlike the molecules you will encounter in yr 10.
Ionic bonding:
Between metals and non-metals.
When metals and non-metals collide, metals will lose their valence electrons to form positive ions (cations), and non-metals will accept these ions, forming negative ions (anions). These ions then form a lattice of alternating ions.
There is an electrostatic force of attraction between unlike ions, however there is a repulsive force between like ions.
The lattice is organised so that there is a greater distance between like ions. As electrostatic forces decrease in strength with an increase in distance, this reduces the repulsive force.
Thus, a net attractive force exists within the ionic lattice. This is called the ionic bond.
Note that the ionic lattice is 3 dimensional.
Metallic bonding:
Only between metals
In metallic solids, metal atoms decide to lose their valence electrons to form positive ions.
*The term used to describe these electrons, which do not orbit any atom and are freely moving, is ==delocalised==.
Thus, this produces an array of metallic cations surrounded by a ==sea of delocalised electrons==.
Note that the 'metallic bond' is the ==electrostatic force of attraction between metal cations and the sea of delocalised electrons==.
Thus, the bond is non-directional.
Covalent Bonding:
Only between non-metals
Non-metal atoms can share electrons to produce a shared pair of electrons.
The covalent bond is the electrostatic force of attraction between a shared pair of electrons produced by 2 atoms, and the 2 atoms' nuclei.
Properties of elements (position $\rightarrow$ properties)
In general, elements closer to the left have a greater metallic nature. Elements closer to the right have a weaker metallic nature.
Metallic nature refers to the readiness of an atom to lose its valence electrons, like how most metals lose their valence electrons to produce the metallic structure.
In terms of chemical reactivity:
For metals, elements become more reactive as they get closer to the bottom right corner of the periodic table.
For non-metals, elements become more reactive as they get closer to the bottom right corner.
This does not include the noble gasses, as they are inert and do not react well.
Nomenclature:
Metals formulae:
Just write the metal's symbol.
e.g. $Ag_{(s)}$
Metals nomenclature:
Just state the metal.
e.g. Silver
Ionic formulae:
Write both metal and non-metal ions, but balance it so that the total charge of the compound is 0.
e.g. $BaCl$ would be wrong, because Barium has a charge of 2+ and Chlorine has a charge of 1-. This would result in a +1 charge.
The correct formulae for Barium chloride is $BaCl_2$.
Positive ions come before negative ions.
Unless the negative ion is $CH_3COO^-$. In this case, we put the positive ion last.
This is because the cation is not attracted to the initial carbon, which has 4 covalent bonds to 3 hydrogens and the other carbon, and thus has an stable electron configuration.
Instead, the last oxygen is actually the atom that is negatively charged. Thus, we place the positive ion last to indicate that it is bonded with the charged oxygen atom, not the carbon atom.
Ionic nomenclature:
Just state the positive and negative ions, in that order.
e.g. Barium Chloride
Do not include charges
e.g. Barium 2+ Chloride - is wrong
UNLESS the positive ion is a transition metal. Then, we have to indicate its valency via roman numerals.
This should be easy to figure out, as you'll be given the ionic compound as a formulae or will be told its valency.
e.g. if we were asked to write the name of $Cu(OH)_2$, from this formula we would be able to deduce Copper's valency in this case to be 2+. Thus, $Cu(OH)_2$'s name would be Copper (II) Hydroxide.
Covalent formulae
Write the elements with the number of elements as subscript.
e.g. Diphosphorus pentoxide is $P_2O_5$.
Covalent nomenclature
For every element, we use a prefix to describe the number of atoms of said element in the given molecule.
The prefixes are (in order of increasing number of atoms (mono = 1)):
Mono
Di
Tri
Tetra
Penta
Hexa
Hepta
Octa
Nona
Deca
HOWEVER, the prefix mono is never used for the first element.
The second element will end with an 'ide' prefix. (root name of the element + ide)
